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Chapter 16: Problem 66

Explain why small, highly charged metal ions are able to undergo hydrolysis.

Short Answer

Expert verified
Small, highly charged metal ions can undergo hydrolysis due to their strong interaction with the polar water molecules (due to their high charge) and the high surface area to volume ratio (due to their small size), which together, significantly increase the probability of hydrolysis.

Step by step solution

01

Understanding Hydrolysis

In hydrolysis, a molecule is split into two parts by the addition of a water molecule. This happens because water molecules, being polar, tend to interact with charged entities, leading to the breaking down of the molecule.
02

Understanding Effective Charge

Highly charged ions, such as small metal ions, have a strong interaction with polar water molecules. The greater the charge, the stronger the interaction, and therefore, the more likely it is for hydrolysis to occur.
03

Understanding Size Effect

Smaller ions have a high ratio of surface area to volume, meaning they have much more surface sites where interactions with water molecules can happen. This increases the probability of hydrolysis happening.
04

Conclusion

Therefore small, highly charged metal ions are more likely to undergo hydrolysis because their high charge and small size allow for a greater interaction with water molecules leading to a higher chance for the breaking of the molecule.

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Most popular questions from this chapter

Calculate the hydrogen ion concentration in moles per liter for each of these solutions: (a) a solution whose \(\mathrm{pH}\) is \(5.20,\) (b) a solution whose \(\mathrm{pH}\) is 16.00 ; (c) a solution whose hydroxide concentration is \(3.7 \times 10^{-9} M\).

In a certain experiment a student finds that the \(\mathrm{pHs}\) of \(0.10 M\) solutions of three potassium salts \(K X, K Y\) and \(\mathrm{KZ}\) are 7.0,9.0 , and 11.0 , respectively. Arrange the acids HX, HY, and HZ in order of increasing acid strength.

Complete this table for a solution: $$\begin{array}{c|c|c}\mathrm{pH} & {\left[\mathrm{H}^{+}\right]} & \text {Solution is } \\\\\hline<7 & & \\\\\hline & <1.0 \times 10^{-7} M & \\\\\hline & & \text { Neutral }\end{array}$$

Like water, ammonia undergoes autoionization in liquid ammonia: $$\mathrm{NH}_{3}+\mathrm{NH}_{3} \rightleftharpoons \mathrm{NH}_{4}^{+}+\mathrm{NH}_{2}^{-}$$ (a) Identify the Bronsted acids and Br酶nsted bases in this reaction. (b) What species correspond to \(\mathrm{H}^{+}\) and \(\mathrm{OH}^{-},\) and what is the condition for a neutral solution?

What are the concentrations of \(\mathrm{HSO}_{4}^{-}, \mathrm{SO}_{4}^{2-},\) and \(\mathrm{H}^{+}\) in a \(0.20 \mathrm{M} \mathrm{KHSO}_{4}\) solution? (Hint: \(\mathrm{H}_{2} \mathrm{SO}_{4}\) is a strong acid: \(K_{\mathrm{a}}\) for \(\left.\mathrm{HSO}_{4}^{-}=1.3 \times 10^{-2} .\right)\)
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